The basic concept of chemical generation and storage of electricity is quite old. Batteries can be traced to 1795 when Volta discovered that two dissimilar metals, placed in an electrically conductive fluid establish an electromotive force from which an electric current can be tapped. Various materials employed as electrodes, a variety of electrolyte possibilities, and numerous combinations of electrodes and electrolytes subsequently have been the object of considerable experimentation for nearly 200 years.
Electrochemical cells generate power when two energetic materials participate in a controlled oxidation-reduction reaction ocurring in an ionically conductive electrolyte. Electrons are transferred in the reaction, and these transferred electrons are collected and distributed by a pair of electrodes generally fabricated from metal, carbon, or the like. Electrons collected at one cell electrode are generally passed through an electrical load circuitry before being returned to the other electrode.
There are two basic electrochemical cell types. Secondary, or rechargeable cells generate electrical energy until some unacceptably low power output is reached caused by depletion of the reactants. Electrical current flow through the cell is thereupon reversed to cause a reversal of the oxidation reduction reaction by which the cell generates power. When a suitable portion of the cell reactants have undergone the reverse reaction, the cell is "recharged" and ready again to provide energy.
Primary cells generate power identically by an oxidation reduction until the reactants contained within the cell have essentially become spent. However, for a variety of reasons having their roots in the nature of the cell reaction or the cell physical configuration, current reversal to recharge the battery is not practical, and the cell is discarded or perhaps recycled.
In much battery development, emphasis has traditionally focused upon locating reactants producing a large amount of energy for a given reactant weight and volume. In addition, it has been necessary to locate conductive electrolytes, chemically compatible with the reactants. However, in larger rechargeable batteries, emphasis has traditionally focused upon improvements to battery electrodes and electrolytes aimed at producing a reasonably efficient battery at very low cost. As a result of this emphasis, these batteries have frequently incorporated low cost individual electrochemical reactants to generate the electromotive energy producing relatively small amounts of energy for a given weight of reactants. As a result of such a low ratio of evolved energy to weight, relatively large amounts of the reactants necessarily have been included in these rechargeable batteries to produce a desired current over a required period of time. For example, according to theoretical calculations, the energy density capability of a lead acid storage battery is about 200 watt-hours per kilogram of reactants.
More recently, in an effort to develop transportation alternates for use in an impending world oil shortage, attempts have been made to power automotive vehicles utilizing electrically powered drive trains drawing electricity from storage batteries contained within the automotive vehicle. An automotive vehicle driven utilizing power provided by batteries carried within the vehicle is transporting the weight of not only the vehicle and its contents, but also of the storage batteries. It is known that vehicle efficiency is strongly dependent upon the weight carried within the vehicle. It has been found that automotive vehicles driven utilizing electrical current from conventional storage batteries having a relatively low energy density generally are not satisfactory. A conventional storage battery providing sufficient electrical current to operate a reasonably commodious automotive vehicle at acceptable speeds and over an acceptable distance is necessarily so weighty that efficient vehicle operation is impaired seriously.
Various attempts have been made to develop a rechargeable storage battery providing a relatively large amount of electromotive energy per unit weight of the battery. Those skilled in the art, referring to the Periodic Table of Elements, have long recognized the alkali and alkaline earth metals and sulfur as possessing the desirable high energy and low weight characteristics. Electrochemical reactions between lower atomic weight alkali metals and sulfur and between lower atomic weight alkaline earth metals and sulfur have long been recognized as potentially providing relatively large energies of reaction from reactants of attractively low weight. For example, according to theoretical calculations, a lithium sulfur battery can produce 2600 watt-hours of energy per kilogram of reactants, a lithium iron disulfide battery about 1100 watt-hours.
A number of proposals have attempted to pair alkali or alkali earth metals with sulfur to produce an efficient storage battery. Many of these proposals have related to primary batteries, that is, batteries designed to use the electrochemical energy of freshly activated battery reactants only once; recharging of these batteries not being contemplated. Alkali or alkaline earth metals reacting with sulfur in such primary batteries have been found to provide acceptable primary battery performance, particularly where an anhydrous electrolyte such as ammonia has been used in the battery. Under anhydrous conditions batteries utilizing, for example, a lithium-sulfur electrochemical reaction pair will produce adequate electrical voltage at operating temperatures significantly below those where an aqueous battery would have become nearly dormant.
Previous proposals for rechargeable batteries utilizing an electrochemical reaction pair involving an alkali or alkaline earth metal and sulfur have proven less satisfactory.
In one proposal, sodium and molten sulfur have been selected as the oxidation reduction reactants. However, the high temperature required in these sodium sulfur batteries has caused serious practical difficulties associated with both heat insulation, particularly of bus bars and bus bar connections, and the make-up supply of heat during periods of extended battery dormancy. Battery housing materials, by necessity, must resist both the elevated temperature and corrosive attack from the reactants. Impurities such as moisture are generally severely dysfunctional to these batteries.
In another high temperature battery, lithium and iron disulfide in a molten salt electrolyte comprise the reactant pair. Beyond the same problems associated with high temperature sodium sulfur batteries, these lithium-iron disulfide cells can suffer from temperature induced iron disulfide instability and short cycle lives attributable at least in part to material migration difficulties.
There have been proposals for the use of an alkali, or alkaline earth metal sulfur, cell-oxidation reduction reaction at ambient temperature. In one such proposal, alkali metal-sulfur electrochemical reaction pairs such as lithium-titanium disulfide or lithium-sulfur have been utilized in combination with organically based electrolytes including dissolved salts. Short cycle lives of such batteries combined with a slow reaction between the organic electrolyte and the alkali metals such as lithium or sodium have dampened development of such batteries. Further, no wholly satisfactory organic electrolyte has yet been found particularly with respect to ionic conductivity. A low ionic conductivity tends not to support adequate cell discharge rates.
In another proposal, electrochemical batteries have utilized a cation producing alkali or alkaline earth metal anode and sulfur cathode together with an electrolyte solute such as an inorganic nitrate or perchlorate of the metal cation dissolved in a cell fluid such as ammonia. Such cells or batteries have demonstrated a capability of being recharged but also have demonstrated rapid decay in cell performance as measured by a reduced current and voltage output of the cell with each subsequent recharging. This decay in cell performance has been attributed to competing reactions between the chemical reactants within the cell, thereby reducing the quantity of chemicals available for storage of current.
It has been further proposed that such cells be divided to separate cell chemical components from one another in an effort to reduce the competing reactions between the electrochemical components. Division has not yet produced a wholly effective rechargeable alkali metal or alkaline earth metal-sulfur battery cell. One significant factor interfering with effective divided cell performance has centered about difficulties in finding a satisfactory divider material that resists destructive effects of fluids in the cell and yet passes metal cations. Another factor has been lack of a suitable supporting electrolyte for use in the cells that does not contribute significantly to competing cell reactions.
Water has long been a favored electrolyte solvent for use in rechargeable batteries. Water has not proven successful as an electrolyte solvent for batteries utilizing alkali and alkaline earth metal-sulfur reactions. Many of the metals react violently with water, and most forms of sulfur useful in the battery reaction are at best insufficiently soluble in the water.
Liquid ammonia exhibits many of the properties that make water a highly desirable battery electrolyte solvent. NH.sub.3 is highly hydrogen bonded, resulting in an unusually elevated boiling point and a substantial heat of vaporization. Ammonia is a protonic, ionizing solvent superior to virtually all but water in dissolving a wide range of electrolyte salts; some salts conduct electricity better in NH.sub.3 than in water.
Yet there are differences between water and liquid ammonia. NH.sub.3 liquid is known to dissolve alkali and alkaline metals to form solutions of a metallic nature when concentrated. These so-called bronzes generally possess characteristics of both electronic and ionic conductance. Such dual properties can be attractive in batteries.
These bronzes have been generally recognized as thermodynamically unstable; some literature reporting half lives as short as 190 hours. Such half lives would preclude utility in most secondary battery applications.
Ammonia is characterized by a theoretical dissociation voltage of 77 millivolts at 25.degree. C. Such a low voltage would seem to seriously limit the use of ammonia in batteries wherein individual cell voltages of in excess of one volt are highly desired. Later evidence has indicated that the actual dissociation voltage of ammonia is significantly in excess of 77 millivolts as a result of significant electrode overvoltages. For these and other reasons, except for use in some primary batteries wherein the ammonia was introduced into the battery immediately prior to battery use, ammonia and ammonia bronzes have not been utilized extensively in batteries.
The use of sulfur as a cathode in conjunction with an ammonia solvent has been suggested since sulfur readily dissolves in ammonia. The kinetics of cell electrochemical reactions in which elemental sulfur is dissolved in ammonia together with supporting electrolytes such as salts have been determined to be quite slow, potentially limiting current flow rates when applied to batteries. Previous proposals have attempted to utilize a sulfur cathode in conjunction with an ammonia electrolyte by the addition of compounds such as alkali metal nitrates, perchlorates, throcyanats, and the like.
It is with cell dividers that additional difficulties with alkali or alkaline earth metal sulfur batteries have been encountered and particularly those with an ammonia electrolyte solvent. It is desirable to separate sulfur from the anode vicinity in such cells to reduce competing cell reactions. Until now, a substantially satisfactory cell separator or partition has not been developed that would (1) retain sulfur in a cell cathode compartment, (2) readily transport cell metal cations, and (3) resist both cell chemical corrosiveness and blocking by-products of competing cell reactions.
In some configurations of a battery using an alkali or an alkaline earth metal anode, it is desirable that anhydrous ammonia be used as an electrolyte within the cell. Often these cells are divided by a partition or a membrane and utilize anhydrous ammonia as an electrolyte. Where a bronze forms between the alkali or alkaline earth metal and the ammonia, this aggressive metal bronze can cause severe damage to, particularly, sensitive cation exchange membranes, but also to partitions such as porous devices. It is known that providing an electrolyte (anolyte) containing a relatively large quantity of salts of the alkali or alkaline earth metal forming the anode in the anode compartment of such a cell can suppress formation of the bronze. However, the relatively large quantities of the salt required for bronze suppression can also prove damaging to performance of the partition or membrane as well by blocking the membrane to ion transport and can promote side reactions in the cell that can markedly decrease performance.